unit 4 review research chem

TYPES OF BONDING

• Ions and ionic charge- Cation: loses electrons, positive charge

  • Anion: gains electrons, negative charge

  • Atoms gain or lose electrons to achieve a stable octet (either 8 valence electrons or 2 for the noble gases in some contexts)

  • Give the ion for the following (example): Ca, S, N, Al

  • In general:

  • Metals (often 1+ or 2+ for main-group metals) tend to form cations

  • Nonmetals tend to form anions

  • Common charges shown in the notes: 1+, 2+, 3+; 1−, 2−, 3−

• Noble gases and ion formation- The most stable/non-reactive chemical family: Noble gases (Group 18)

  • Atoms gain or lose electrons to obtain the same electron configuration as the nearest noble gas (based on atomic number)

• Bohr models (brief idea)- Model showing electrons in defined shells around the nucleus

  • Used here to illustrate electron configurations and ion formation (e.g., Na vs Ne vs H vs F)

• Law of conservation of mass and electrons- Mass/matter cannot be created or destroyed in reactions

  • Electrons are involved in bonding and can be gained/lost (ions) or shared (bonds)

• Chemical bonding overview- Compounds form when atoms combine; bonds hold atoms together

  • Bonds arise from interactions between nuclei and valence electrons of other atoms

  • The phrase often repeated:

  • “chemical bonds” = the force of attraction between nuclei and valence electrons of other atoms

IONIC COMPOUNDS

• Ionic bonds- Formed by transfer of valence electrons from a metal to a non-metal

  • Resulting bond is between oppositely charged ions (cation and anion)

  • Most common combination: Metal + Non-metal (e.g., NaCl; Na has 1 valence electron, Cl has 7)

  • Example ion formation: Na → Na⁺ (loses 1 e⁻), Cl → Cl⁻ (gains 1 e⁻)

• Ionic compounds and crystal lattice- Atoms form an ionic lattice with alternating cations and anions

  • Lattice structure: repeating pattern of opposite charges

• Properties of ionic compounds- Rigid and brittle

  • High melting points

  • Conduct electricity when melted or dissolved in water (ionic compounds are electrolytes when molten or in solution)

  • Often referred to as salts; high solubility in water in many cases

• Examples and visuals from the notes- Na and Cl with valence electrons: Na (1 Ve⁻), Cl (7 Ve⁻) → Na⁺ and Cl⁻ after electron transfer

  • The shift to “O” valence electrons in different slide visuals shows the concept of achieving full valence shells (8 valence electrons for most, 2 for H/He)

COVALENT BONDS

• Covalent bonds- Formed by sharing valence electrons between non-metals

  • Electrons are attracted to the nuclei of both atoms; neither nucleus gains full control (bonding pair)

  • Generally occurs in non-metal + non-metal combinations

• Shared electrons and octet rule- Atoms share electrons to achieve a noble gas configuration (stable octet, 8 valence electrons) or duet for H/He

• Properties of covalent (molecular) compounds- Do not conduct electricity in solid state

  • Low melting/boiling points compared to ionic compounds

  • Weaker than ionic solids

• Notable terms- Covalent bond: a bond formed from sharing of valence electrons

  • Molecule: a particle in which atoms are joined by covalent bonds

  • Molecular compound: a pure substance formed from multiple molecules

POLYATOMIC IONS AND MOLECULES

• Polyatomic ions- An ion made up of more than one atom that acts as a single unit

  • Examples:

  • 
    

    Polyatomic ion = an ion made up of more than one atom, which acts as one single unit

  • In Lewis structures, count all valence electrons contributed by all atoms, then adjust for the charge of the ion

• Diatomic elements- Elements that commonly exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂ (and sometimes others under certain conditions)

  • In the slides, diatomic elements list is shown (e.g., Br, I, N, Cl, H, O, F)

• Polyatomic ions and the formal charge concept (in Lewis structures)- When drawing Lewis structures for ions like ClO⁻, add the extra electron due to the negative charge

  • Total valence electrons = sum of valence electrons of each atom + any charge adjustment

LEWIS STRUCTURES

• What they are- Visual representation of covalent molecules and covalent bonds

• Octet rule, duet rule- Most elements seek 8 valence electrons (octet)

  • Hydrogen and Helium are satisfied with 2 valence electrons (duet)

• Steps to draw Lewis structures 1) Determine the total number of valence electrons in the molecule (TotalValence = sum of valence electrons of all atoms) - Example: CH₄: C = 4 valence e⁻, H = 1 valence e⁻ each, total = 4 + 4×1 = 8 valence electrons

  • Symbolically: $V_{\text{total}} = \text{sum of valence electrons of all atoms}$
    

    2) Determine the central atom

  • Rules (as given): Always carbon as a central in simple organics; Never hydrogen; If multiple possibilities, choose the element that there is less of; choose the least electronegative element
    

    3) Place central atom and surrounding terminal atoms

  • Total valence electrons are arranged with central atom in the center and terminal atoms around it
    

    4) Draw bonds to connect terminal atoms to the central atom

  • Each bond represents 2 electrons (a bonding pair)

  • 1 bond = 2 electrons

  • Example: CH₄: draw four C–H single bonds
    

    4a) Subtract electrons used for bonds from total valence electrons

  • Total valence e⁻ minus electrons used in bonds (e.g., 4 bonds × 2 e⁻ = 8 e⁻ for CH₄)
    

    4b) Add lone pairs to terminal atoms first to satisfy their octets

  • Lone pair = 2 electrons
    

    4c) Stop adding lone pairs when an atom reaches its full valence shell (8 for most elements, 2 for H/He)
    

    4d) If electrons remain after all terminal atoms have octets, place the remaining lone pairs on the central atom
    

    5) Check your structure

  • All atoms (except H/He) should have a full octet; ensure the total number of electrons matches the original count

• Examples from the notes- Methane, CH₄: central C, four H around; 4 bonds, 8 valence electrons used; each H has 2 electrons around it; carbon ends with 8 electrons total

  • NF₃: Total valence e⁻ = N(5) + 3×F(7) = 5 + 21 = 26; central atom N; draw bonds to three F; complete octets on F; place remaining electrons on N as needed; formal charges considered later

• Special cases and checks- Boron and other group 13 elements: can be satisfied with 6 electrons (expanded discussion in notes)

  • Expanded octet for sulfur, phosphorus, etc. (more than 8 valence electrons)

• Formal charges (for validating Lewis structures)- Formal charge on atom = (#valence electrons) − [(bonding electrons) + (nonbonding electrons)]

  • Best structures have as many zero formal charges as possible; when not possible, place negative charges on more electronegative atoms

  • Example approach: N–O vs N–O resonance structures; choose resonance that minimizes formal charges and places negative charges on more electronegative atoms

• Resonance structures- When more than one correct Lewis structure exists, the actual structure is a resonance hybrid

  • Represented as resonance structures with double-headed arrows

EXAMPLE LEWIS STRUCTURE PRACTICE

• CH₄ (methane)- Total valence electrons: $4 + 4 \times1 = 8$

  • Central atom: C; terminal hydrogens around; 4 single bonds; no lone pairs on C; octet satisfied for all atoms

• NF₃ (nitrogen trifluoride)- Total valence electrons: $N(5) + 3 \times F(7) = 26$

  • Central atom: N; 3 F around; draw 3 bonds; complete F octets; remaining electrons placed as lone pairs on N if needed; then consider formal charges

• Oxygens in ClO⁻ (chlorite ion) as an example of a polyatomic ion Lewis structure- Total valence electrons: Cl(7) + O(6) + extra electron for −1 charge = 14 valence electrons total

  • Follow steps 1–4 to produce the best resonance structure; adjust for the negative charge by adding one electron initially

• Formal charge quick reference- For a given structure, compute FC for each atom and aim to minimize negative charges on less electronegative atoms and place negatives on more electronegative atoms when possible

DIATOMIC MOLECULES, POLYATOMIC IONS, AND EXCEPTIONS TO OCTET

• Diatomic elements- Br, I, N, Cl, H, O, F; common diatomic molecules include H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂

• Exceptions to the octet rule- Group 13 elements (e.g., Boron) may be satisfied with 6 electrons in simple compounds

  • Expanded octet for elements like sulfur and phosphorus can exceed 8 valence electrons

MOLECULAR SHAPES (VSEPR)

• Key idea- Shapes depend on the repulsion between electron domains (bonding and lone pairs) around a central atom

  • Lone pairs repel more strongly than bonding pairs, affecting bond angles

• Core vocabulary- Steric number / Electron domain: number of regions of electron density around the central atom (bonding pairs + lone pairs)

  • Electron geometry: arrangement of electron domains around the central atom

  • Molecular geometry: arrangement of atoms around the central atom

• VSEPR model steps (example CH₄)- Step 1: Determine electron domains = 4 (4 bonding pairs, no lone pairs) → Electron geometry: tetrahedral

  • Step 2: Central atom: C; terminal atoms around

  • Step 3: Place atoms around central atom

  • Step 4: Determine bond angles from electron geometry (for tetrahedral, ~109.5°)

• Examples from notes- CH₄: Electron geometry = tetrahedral; Molecular geometry = tetrahedral; Bond angle ≈ 109.5°

  • NH₃ (ammonia): Electron geometry = tetrahedral; Molecular geometry = trigonal pyramidal; Bond angle ≈ 107° (slightly less than 109.5° due to lone pair repulsion)

  • NF₃: Similar to NH₃ in geometry (pyramidal) with bond angles slightly less than 109.5°

  • CO₂: Electron geometry = linear; Molecular geometry = linear; Bond angle = 180°

  • H₂O: Electron geometry = tetrahedral; Molecular geometry = bent; Bond angle ≈ 104.5°

• Multi-centered molecules- For molecules with multiple central atoms, treat one center at a time for the purpose of predicting geometry; compex molecules may require considering multiple local centers

POLARITY AND INTERMOLECULAR FORCES (IMFs)

• Bond polarity- Non-polar covalent bonds: electrons shared equally between identical or very similar atoms; no partial charges

  • Polar covalent bonds: electrons shared unequally; results in partial charges (δ+ and δ−) on atoms

  • Dipole moment is a vector quantity indicating polarity; molecules with polar bonds can be non-polar overall if symmetry cancels dipoles

• Polar vs non-polar molecules- A molecule is polar if it has at least one polar bond and the dipoles do not cancel due to asymmetrical geometry

  • A molecule is non-polar if it is symmetrical and dipoles cancel out (e.g., CO₂ is linear and non-polar, CH₄ is symmetrical and non-polar despite C–H bonds being slightly polar)

• Example visuals- H–Cl has a polar covalent bond; Cl is more electronegative than H, causing δ− on Cl and δ+ on H

  • C=O bonds are individually polar, but CO₂ is non-polar due to symmetrical linear geometry

• Molecular polarity and symmetry- Symmetrical molecules tend to be non-polar; asymmetrical molecules with lone pairs or different terminal atoms tend to be polar

• Intermolecular forces (IMFs)- Intermolecular forces are forces between molecules, not within a molecule

  • Types of IMFs (from weakest to strongest, as noted):

  • London Dispersion Forces (LDFs) — present in all molecules, especially non-polar

  • Dipole–dipole interactions — occur between polar molecules; stronger than LDFs

  • Hydrogen bonding — a special, particularly strong dipole-dipole interaction; occurs when H is bonded to F, O, or N

  • Hydrogen bonding explained with F, O, N and a hydrogen attached to them

• Molecular symmetry and polarity visuals- Symmetrical examples (non-polar): many hydrocarbons (H–C–H frameworks with only C and H)

  • Asymmetrical examples (polar): molecules with lone pairs or different substituents on the central atom

SUMMARY OF BONDING (COMPARISON)

• Ionic bonds- Transfer of electrons from cation to anion

  • Bond formation leads to crystal lattices

  • Structure: crystal lattice; Anions and cations arrange in a repeating pattern

  • Physical properties: solid at room temperature, high melting points

  • Electrical properties: good conductors when melted or dissolved in water

  • Typical structure: metal + non-metal (e.g., NaCl)

• Covalent bonds- Sharing of electrons between two non-metals

  • Resulting structures: molecules or networks

  • General properties: poor electrical conductivity in solids, variable melting points, weaker than ionic solids in many cases

• Metallic bonds- Delocalized electrons (sea of electrons) shared among metal atoms

  • Properties: good electrical and thermal conductivity, malleable, ductile

• Summary table (as per slides)- Bond formation: Ionic (transfer), Covalent (sharing), Metallic (delocalized electrons)

  • Type of structure: Crystal lattice (ionic), Molecular shapes (covalent), “Electron sea” (metallic)

  • Physical state: Solid (ionic and many covalent solids), Liquid/Gas (some covalent networks), Solid (metals, etc.)

  • Melting point: High (ionic), Low (molecular covalent), Very high (metals in many cases)

  • Solubility in water: Good (ionic), Poor/none (many covalent), varies (metals)

  • Conductivity: Good (ionic when molten/dissolved), Poor/none (molecular), Good (metals when solid and molten)

  • Other properties: Brittleness (ionic), Malleable (metallic)

INTERMOLECULAR FORCES (DEEPER)

• What are IMFs?- Forces that hold multiple molecules together in condensed phases

  • They are weaker than intramolecular (chemical) bonds

• Key IMF types covered- London dispersion forces (LDFs) — present in all molecules; especially important for non-polar species

  • Dipole–dipole interactions — between polar molecules; stronger than LDFs

  • Hydrogen bonds — strongest among the IMFs discussed; occurs when hydrogen is bonded to F, O, or N

• Practical implications- IMFs determine boiling/melting points, viscosity, surface tension, and miscibility

  • Non-polar molecules typically have lower MP/BP due to reliance on LDFs

MISCELLANEOUS TOPICS LINKING THEM TO REAL WORLD/FOUNDATIONS

• Connections to foundational chemistry- Octet rule and noble gas configurations underpin why atoms form certain bonds and charges

  • The concept of formal charges helps evaluate the most plausible Lewis structures, including resonance structures

  • VSEPR links electron-domain geometry to molecular shapes, which in turn affect polarity and IMFs

• Relevance to materials and real-world substances- Ionic compounds form salts used in metallurgy and chemistry; their properties underpin many applications in energy, environment, and biology

  • Covalent compounds include many organic molecules essential to life and everyday materials; intermolecular forces influence properties like boiling point and solubility

  • Metallic bonding explains properties of metals used in construction, electronics, and industrial processes

KEY FORMULAS AND NUMERICAL REMINDERS (IN LAtex)

• Octet vs duet rule

  • Most atoms want 8 valence electrons; hydrogen and helium want 2 valence electrons: \text{valence target} = \begin{cases} 8, & \text{for most elements}\ 2, & \text{for H, He} \ \text{(duet rule)} \

  \text{ }\n
  \text{ }\n

\exists \n

\text{example}\n

\text{o}\n

\text{d}\n

\end{cases}

• Total valence electrons in a molecule: $V_{\text{total}} = \text{sum of valence electrons of all atoms}$

• 1 bond equals 2 electrons: $1 \text{ bond} = 2 \text{ electrons}$

• Formal charge on an atom: $FC = V - (B + N)$

• Where V = number of valence electrons for the atom, B = number of electrons in bonds (counted as bonding electrons per atom? typically half the bond electrons per atom), and N = number of nonbonding electrons on the atom

• For resonance and structure validation: maximize zero formal charges when possible; place negative charges on more electronegative atoms

• Bond angles (typical values from VSEPR):- Tetrahedral (CH₄): $$\theta \approx 10